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Group 1 and group 2: Alkali and alkaline earth metals

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Intros
Lessons
  1. Group 1 and group 2: a summary
  2. Intro to group 1 and 2.
  3. Properties of group 1 and group 2 metals.
  4. How do the properties change going down the group? (Trends in properties)
  5. Why do group 1 and 2 metals get more reactive down the group?
  6. Reactions of group 1 and 2 metals.
  7. Solubility of group 1 and 2 metals.
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Examples
Lessons
  1. Recall the trend in properties of the alkali metals.
    Sort the alkali metals for the following properties:
    i) From hardest to softest.
    ii) From lowest to highest melting point.
    1. Recall and predict the reactions of the alkali metals.
      In a school laboratory, a chunk of sodium metal is put in water.

      i) Write a full word equation to show the reaction taking place.

      ii) One of the products of the reaction can be identified by using universal indicator. What color will the solution become if universal indicator is added?

      iii) After this, a student puts a chunk of potassium metal in a separate bath of water. Will this reaction be more or less vigorous? Explain why.
      Topic Notes
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      Introduction

      Alkali metals and alkaline earth metals, found in Groups 1 and 2 of the periodic table elements, respectively, are essential elements in chemistry. The introduction video provides a comprehensive overview of these elements, offering valuable insights into their unique properties and behaviors. Alkali metals, including lithium, sodium, and potassium, are highly reactive and soft, while alkaline earth metals, such as beryllium, magnesium, and calcium, exhibit similar but less extreme characteristics. These elements play crucial roles in various chemical reactions and have numerous applications in industry and everyday life. Their distinctive properties, such as low ionization energies and high reactivity with water, make them fascinating subjects of study. Understanding the behavior of alkali and alkaline earth metals is fundamental to grasping broader concepts in chemistry and the periodic table elements. The video serves as an excellent starting point for exploring these important elements and their significance in the field of chemistry.

      Properties of Alkali and Alkaline Earth Metals

      Alkali metals and alkaline earth metals are two distinct groups of elements in the periodic table, each with unique physical and chemical properties. Understanding these properties is crucial for comprehending their behavior and applications in various fields. This article will explore and compare the characteristics of these two important metal groups.

      Alkali metals, found in Group 1 of the periodic table, include elements such as lithium, sodium, and potassium. These metals are known for their exceptional reactivity, softness, low density, and low melting points. The physical properties of alkali metals make them stand out among other elements. They are silvery-white in appearance and extremely soft, easily cut with a knife. Their low density allows them to float on water, a unique characteristic among metals.

      The chemical properties of alkali metals are equally fascinating. They readily lose their outermost electron, making them highly reactive. This reactivity is so intense that they must be stored in oil to prevent reaction with air or water. When exposed to water, alkali metals react vigorously, often producing hydrogen gas and heat. This reaction becomes more violent as we move down the group, with cesium and rubidium reacting explosively.

      Alkaline earth metals, found in Group 2, include elements like beryllium, magnesium, and calcium. While sharing some similarities with alkali metals, they have distinct properties of their own. The physical properties of alkaline earth metals show some parallels to alkali metals, such as being silvery-white and relatively soft, though not as soft as their Group 1 counterparts. They also have low densities, but higher than those of alkali metals.

      The chemical properties of alkaline earth metals reveal both similarities and differences compared to alkali metals. Like alkali metals, they are reactive, but generally less so. This reduced reactivity is due to their electronic configuration, having two valence electrons instead of one. Alkaline earth metals still react with water, but the reaction is typically less vigorous than that of alkali metals. For instance, calcium reacts with water to produce calcium hydroxide and hydrogen gas, but the reaction is not as explosive as sodium's reaction with water.

      Comparing the two groups, we find that alkali metals are generally more reactive than alkaline earth metals. This is evident in their reaction with oxygen; alkali metals form oxides, peroxides, and superoxides, while alkaline earth metals typically form only oxides. Both groups have low melting points compared to transition metals, but alkali metals have even lower melting points than alkaline earth metals.

      The softness of these metals is another point of comparison. Alkali metals are exceptionally soft and can be cut with a knife, while alkaline earth metals, though softer than many other metals, are harder than alkali metals. This difference in hardness is related to their atomic structure and bonding.

      Both groups share the property of low density, but alkali metals take this to an extreme. Lithium, sodium, and potassium can all float on water, a property not shared by alkaline earth metals. This low density is due to their large atomic radii and relatively few electrons.

      The low melting points of both groups are noteworthy, especially for alkali metals. Cesium, for example, melts at just 28.5°C (83.3°F), barely above room temperature. Alkaline earth metals have higher melting points than alkali metals but still lower than many other elements.

      In terms of practical applications, the properties of these metals make them valuable in various industries. Sodium is used in street lamps and industrial processes, while lithium is crucial in battery technology. Magnesium, an alkaline earth metal, is used in lightweight alloys for aerospace and automotive industries.

      Understanding the properties of alkali and alkaline earth metals is essential in chemistry and materials science. Their unique characteristics, from high reactivity to low density and melting points, make them fascinating subjects of study and valuable resources in technology and industry. As we continue to explore and utilize these elements, their importance in our daily lives and technological advancements becomes increasingly apparent.

      Electronic Structure and Reactivity Trends

      The electronic structure of alkali and alkaline earth metals plays a crucial role in determining their chemical properties and reactivity. These elements, found in Groups 1 and 2 of the periodic table, respectively, have unique characteristics due to their outer shell electron configurations. Understanding their electronic structure is key to comprehending their behavior in chemical reactions and their trends in reactivity.

      Alkali metals, located in Group 1, have a single electron in their outermost shell. This lone valence electron is easily lost, resulting in the formation of a +1 ion. The electronic configuration of alkali metals can be represented as [noble gas]ns¹, where 'n' is the principal quantum number of the outermost shell. For example, sodium (Na) has the electronic configuration [Ne]3s¹.

      Alkaline earth metals, found in Group 2, have two electrons in their outermost shell. Their electronic configuration can be written as [noble gas]ns². For instance, magnesium (Mg) has the configuration [Ne]3s². These two outer shell electrons can be removed to form +2 ions, albeit with more energy required compared to alkali metals.

      The number of outer shell electrons significantly affects the properties of these elements. With fewer electrons in the outermost shell, these metals have lower ionization energies compared to other elements in the same period. This makes them highly reactive, as they readily lose electrons to form positive ions.

      As we move down the groups of alkali and alkaline earth metals, a clear trend in reactivity emerges. The reactivity increases as we go from top to bottom in both groups. This trend is directly related to the atomic size and the concept of electron shielding.

      Electron shielding, also known as screening effect, occurs when inner shell electrons shield the outer electrons from the full attractive force of the nucleus. As we move down a group, the number of inner electron shells increases, leading to greater shielding. This results in a weaker attraction between the nucleus and the outermost electrons.

      The impact of electron shielding on reactivity is significant. With increased shielding, the valence electrons are held less tightly by the nucleus, making them easier to remove. This leads to lower ionization energies and higher reactivity as we move down the group.

      For Group 1 elements (alkali metals), the reactivity trend is particularly pronounced. Lithium (Li) at the top is the least reactive, while francium (Fr) at the bottom is the most reactive. This is evident in their reactions with water; while lithium reacts slowly, elements like sodium and potassium react vigorously, and rubidium and cesium react explosively.

      Similarly, in Group 2 (alkaline earth metals), the reactivity increases from beryllium (Be) to radium (Ra). Beryllium does not react with water, while magnesium reacts slowly. Calcium and the elements below it react more readily with water, forming hydroxides and releasing hydrogen gas.

      The trend in ionization energy further supports this reactivity pattern. As we move down the group, the ionization energy decreases due to increased atomic size and electron shielding. This makes it easier for the atoms to lose electrons and form positive ions, enhancing their reactivity.

      It's important to note that while the general trend shows increasing reactivity down the group, there can be some exceptions or variations due to other factors such as lattice energy and hydration energy of the resulting ions.

      In conclusion, the electronic structure of alkali and alkaline earth metals, characterized by their outer shell electrons, plays a pivotal role in determining their chemical behavior. The trend of increasing reactivity as we move down these groups is a direct consequence of increasing atomic size and electron shielding. This understanding of electronic structure and reactivity trends is fundamental in predicting and explaining the chemical properties and reactions of these important groups of elements in the periodic table.

      Common Reactions of Group 1 and Group 2 Elements

      Alkali metals (Group 1) and alkaline earth metals (Group 2) are known for their high reactivity due to their electron configurations. These elements readily participate in various chemical reactions, particularly with water, oxygen, and chlorine. Understanding these reactions is crucial for comprehending the behavior of these elements in different environments.

      Reactions with Water:

      Alkali metals react vigorously with water, forming hydroxides and releasing hydrogen gas. The general equation for this reaction is:

      2M + 2HO 2MOH + H (where M represents an alkali metal)

      For example, sodium reacts with water as follows:

      2Na + 2HO 2NaOH + H

      This reaction is highly exothermic, often resulting in the ignition of the hydrogen gas produced. The intensity of the reaction increases down the group, with cesium reacting explosively with water.

      Alkaline earth metals also react with water, but less vigorously compared to alkali metals. The general equation is:

      M + 2HO M(OH) + H (where M represents an alkaline earth metal)

      For instance, calcium reacts with water as follows:

      Ca + 2HO Ca(OH) + H

      The reactivity of alkaline earth metals with water increases down the group, with beryllium being the least reactive and radium the most reactive.

      Reactions with Oxygen:

      Alkali metals react rapidly with oxygen in the air, forming oxides. The general equation is:

      4M + O 2MO (where M represents an alkali metal)

      For example, lithium reacts with oxygen to form lithium oxide:

      4Li + O 2LiO

      Sodium and potassium form peroxides (NaO and KO) when reacting with oxygen, while rubidium and cesium form superoxides (RbO and CsO).

      Alkaline earth metals also react with oxygen, forming oxides. The general equation is:

      2M + O 2MO (where M represents an alkaline earth metal)

      For instance, magnesium burns in oxygen to form magnesium oxide:

      2Mg + O 2MgO

      This reaction is often demonstrated in chemistry classes due to the bright white light produced during the combustion.

      Reactions with Chlorine:

      Both alkali and alkaline earth metals react readily with chlorine gas to form chlorides. For alkali metals, the general equation is:

      2M + Cl 2MCl (where M represents an alkali metal)

      For example, potassium reacts with chlorine to form potassium chloride:

      2K + Cl 2KCl

      Alkaline earth metals react similarly, but form dichlorides:

      M + Cl MCl (where M represents an alkaline earth metal)

      For instance, calcium reacts with chlorine to form calcium chloride:

      Ca + Cl CaCl

      These reactions are highly exothermic and often produce a bright flame. The resulting chlorides are ionic compounds with high melting and boiling points.

      Observed Phenomena:

      During these reactions, several interesting phenomena

      Flame Tests and Spectral Lines

      Flame tests are a fundamental analytical technique used to identify alkali and alkaline earth metals in chemistry. This method relies on the characteristic colors produced when these elements are heated in a flame, providing a quick and simple way to determine their presence. The concept of flame tests is closely tied to the unique electron configurations of these elements and their behavior when exposed to high temperatures.

      When an alkali or alkaline earth metal is introduced into a flame, the heat causes the atoms to become excited. As electrons absorb energy, they jump to higher energy levels. When these excited electrons return to their ground state, they release energy in the form of light, producing distinctive spectral lines and colors. Each element has a unique electron configuration, resulting in specific wavelengths of light being emitted, which we perceive as different colors.

      The characteristic colors produced by common elements in flame tests are as follows:

      • Lithium: Bright red
      • Sodium: Intense yellow
      • Potassium: Pale lilac or violet
      • Calcium: Brick red
      • Strontium: Crimson red
      • Barium: Pale green
      • Copper: Blue-green

      The relationship between flame colors and electron transitions is crucial to understanding the underlying principles of flame tests. The color observed corresponds to the energy difference between the excited state and the ground state of the electrons. This energy difference is unique to each element, which is why different elements produce distinct colors. The spectral lines produced in flame tests are a direct result of these specific electron transitions.

      Flame tests play a significant role in qualitative analysis and element identification. They offer a quick, cost-effective method for detecting the presence of certain metals in a sample. This technique is particularly useful in fields such as forensic science, environmental monitoring, and quality control in manufacturing processes. While flame tests are not as precise as more advanced spectroscopic methods, they remain a valuable tool for preliminary identification and educational purposes.

      In conclusion, flame tests provide a visually striking and scientifically sound method for identifying alkali and alkaline earth metals. By understanding the connection between flame colors, spectral lines, and electron transitions, chemists can quickly and effectively determine the presence of specific elements in a sample. This technique continues to be an essential part of chemical analysis and serves as an excellent introduction to the principles of atomic structure and energy levels.

      Solubility Trends in Group 1 and Group 2 Compounds

      Understanding the solubility trends of Group 1 and Group 2 compounds is crucial in chemistry, as it provides insights into their behavior in aqueous solutions. These trends are particularly evident when examining the solubility of hydroxides, sulfates, and carbonates across both groups. The general pattern shows increasing solubility down Group 1 (alkali metals) and decreasing solubility down Group 2 (alkaline earth metals) for certain compounds, though there are notable exceptions.

      Group 1 hydroxides exhibit a clear trend of increasing solubility as we move down the group. Lithium hydroxide (LiOH) is the least soluble, while cesium hydroxide (CsOH) is the most soluble. This trend is attributed to the decreasing lattice enthalpy and increasing hydration enthalpy of the larger cations as we descend the group. The larger cations have weaker attractions to the hydroxide ions, making it easier for water molecules to separate them.

      In contrast, Group 2 hydroxides show a reverse trend. Beryllium hydroxide (Be(OH)2) is slightly soluble, while magnesium hydroxide (Mg(OH)2) is sparingly soluble. However, as we move down to calcium hydroxide (Ca(OH)2), strontium hydroxide (Sr(OH)2), and barium hydroxide (Ba(OH)2), solubility increases. This anomaly is due to the high charge density of Be2+ and Mg2+ ions, which leads to stronger attractions with hydroxide ions.

      When it comes to sulfates, Group 1 compounds are generally highly soluble in water. From lithium sulfate (Li2SO4) to cesium sulfate (Cs2SO4), all are readily soluble. This high solubility is due to the strong hydration of both the Group 1 cations and the sulfate anion. In Group 2, however, we observe a decrease in solubility as we move down the group. Beryllium sulfate (BeSO4) and magnesium sulfate (MgSO4) are highly soluble, but calcium sulfate (CaSO4) has low solubility. Strontium sulfate (SrSO4) and barium sulfate (BaSO4) are practically insoluble. This trend is explained by the increasing size of the Group 2 cations, which leads to stronger lattice energies that outweigh the hydration energies.

      Carbonates present an interesting case. In Group 1, all carbonates are soluble, with solubility increasing down the group. Lithium carbonate (Li2CO3) is the least soluble, while cesium carbonate (Cs2CO3) is the most soluble. This trend aligns with the general pattern observed in Group 1 compounds. However, Group 2 carbonates show very low solubility across the board. From beryllium carbonate (BeCO3) to barium carbonate (BaCO3), all are practically insoluble in water. This insolubility is due to the high lattice energies of these compounds, which significantly exceed their hydration energies.

      It's important to note some exceptions to these general trends. For instance, lithium compounds often behave differently from other Group 1 elements due to the small size and high charge density of the Li+ ion. Lithium carbonate, for example, has a much lower solubility compared to other Group 1 carbonates. Similarly, beryllium compounds in Group 2 often show anomalous behavior due to the small size and high polarizing power of the Be2+ ion.

      Understanding these solubility trends is essential in various applications, from water treatment and mineral processing to the development of pharmaceuticals. The solubility of hydroxides is particularly relevant in pH control and precipitation reactions, while the low solubility of certain sulfates and carbonates is exploited in analytical chemistry for separation and identification of ions. These trends also play a crucial role in geochemistry, influencing the formation and dissolution of minerals in natural water systems.

      In conclusion, the solubility trends of Group 1 and Group 2 compounds provide a fascinating insight into the interplay

      Applications and Importance in Daily Life

      Alkali and alkaline earth metals play crucial roles in our daily lives, with applications spanning various industries and technologies. These elements are fundamental to many products we use every day, from electronics to construction materials and medical treatments. The applications of alkali metals, such as sodium, potassium, and lithium, are particularly noteworthy in modern technology and industrial processes. Sodium, for instance, is essential in the production of street lamps, providing efficient lighting in urban areas. Lithium has revolutionized the electronics industry, becoming a key component in rechargeable batteries for smartphones, laptops, and electric vehicles. Potassium finds extensive use in fertilizers, enhancing crop yields and supporting global food production. The applications of alkaline earth metals are equally diverse and significant. Calcium, a prominent alkaline earth metal, is vital in the construction industry, forming the basis of cement and concrete. It's also crucial in the human body, supporting bone health and muscle function. Magnesium, another alkaline earth metal, is widely used in the automotive and aerospace industries due to its lightweight properties and strength. In everyday products, these metals are often hidden yet indispensable. Beryllium, for example, is used in X-ray machines and computer components, contributing to medical diagnostics and information technology. Strontium finds application in fireworks, adding vibrant red colors to displays. In the realm of technological applications, barium is used in CRT screens and spark plugs, while radium was historically used in luminous paints for watch dials. The medical field also benefits greatly from these elements. Lithium compounds are used in the treatment of bipolar disorder, while magnesium sulfate is administered to prevent seizures in pregnant women with preeclampsia. Calcium-based medications are common in treating osteoporosis and calcium deficiencies. The industrial uses of alkali and alkaline earth metals extend to water treatment, where sodium and calcium compounds are used to soften hard water. In the chemical industry, these metals serve as catalysts in various reactions, facilitating the production of plastics, pharmaceuticals, and other essential products. Their reactivity makes them valuable in neutralizing acids, a property utilized in environmental clean-up efforts and industrial waste management. As we continue to advance technologically, the importance of alkali and alkaline earth metals in our daily lives is likely to grow, driving innovation in energy storage, materials science, and environmental technologies.

      Conclusion

      Alkali and alkaline earth metals play crucial roles in chemistry and everyday life. Alkali metals, found in Group 1 of the periodic table, are highly reactive and soft, with low melting points. They readily lose their outermost electron, forming positive ions. Alkaline earth metals, located in Group 2, are slightly less reactive but still form positive ions easily. Both groups share similarities in their chemical behavior, such as reacting vigorously with water and forming strong bases. Understanding their properties is essential for various applications, from industrial processes to biological systems. Alkali metals are vital in batteries and nuclear reactors, while alkaline earth metals are crucial in construction materials and medical treatments. Their unique characteristics, including high conductivity and low density, make them indispensable in modern technology. Mastering the properties of these elements is fundamental for students and professionals in chemistry, engineering, and related fields, enabling innovation and problem-solving in numerous scientific and industrial applications.

      Group 1 and Group 2: A Summary

      Group 1 and Group 2: A Summary
      Introduction to Group 1 and Group 2.

      Step 1: Introduction to Group 1 and Group 2 Elements

      Group 1 and Group 2 elements are known as the alkali and alkaline earth metals, respectively. These groups are characterized by their high reactivity, which distinguishes them from other metals. Group 1 elements have one outer shell electron, while Group 2 elements have two outer shell electrons. This similarity in their electronic structure results in similar properties between the two groups.

      Step 2: Properties of Alkali and Alkaline Earth Metals

      Both groups exhibit certain properties that change as you move down the group. These properties are influenced by the electronic structure, particularly the outer shell electrons and the number of inner shells. The properties of these metals are determined by their outer shell electrons. Group 1 elements are highly reactive due to their single outer shell electron, while Group 2 elements, with two outer shell electrons, are also quite reactive but to a slightly lesser extent.

      Step 3: Common Reactions of Group 1 and Group 2 Elements

      Group 1 and Group 2 elements are known for their reactivity, especially with oxygen and water. This reactivity has made them well-studied and known for a long time. Historically, these elements were first identified in their compound forms rather than as pure metals. For example, magnesium oxide (magnesia) and calcium oxide (lime) were known compounds long before the pure metals were isolated.

      Step 4: Reactivity and Historical Context

      The reactivity of these metals with oxygen and water means they are often found in nature as compounds rather than in their pure elemental form. Chemists initially knew these elements through their oxides, which are salts composed of a metal positive ion and an oxide negative ion. These oxides produce alkaline solutions when dissolved in water, leading to the names "alkali" and "alkaline earth" metals.

      Step 5: Solubility Differences Between Group 1 and Group 2

      One key difference between Group 1 and Group 2 elements is their solubility in water. Group 1 compounds are generally quite soluble, while Group 2 compounds are less so. Despite this difference, both groups are highly reactive and were originally found as compounds before being isolated as elements. This reactivity is due to the presence of oxygen in the air, which readily reacts with these metals.

      Step 6: Preservation and Practical Applications

      Due to their reactivity, small samples of Group 1 and Group 2 metals are often preserved in oil to prevent oxidation. When exposed to air, these metals quickly form oxides on their surfaces. This reactivity has allowed chemists to conduct extensive research on these elements, leading to a deep understanding of their properties and behaviors.

      Step 7: Conclusion and Further Exploration

      In summary, Group 1 and Group 2 elements are characterized by their high reactivity, similar electronic structures, and historical significance. Their interactions with oxygen and water have made them well-known and extensively studied. Understanding these properties provides a foundation for further exploration of their chemical behaviors and practical applications.

      FAQs

      Here are some frequently asked questions about alkali and alkaline earth metals:

      What are the 5 properties of alkaline earth metals?

      The five key properties of alkaline earth metals are: 1) They are silvery-white in color, 2) They have relatively low densities, 3) They are good conductors of heat and electricity, 4) They have relatively high melting and boiling points compared to alkali metals, and 5) They are highly reactive, though less so than alkali metals.

      Are alkaline earth metals soft or hard?

      Alkaline earth metals are generally softer than most other metals but harder than alkali metals. Their hardness increases as you move down the group in the periodic table. For example, beryllium is relatively hard, while barium is quite soft.

      What are the main characteristics of alkali metals?

      The main characteristics of alkali metals include: 1) They are extremely reactive, 2) They have low melting and boiling points, 3) They are soft and can be cut with a knife, 4) They have low densities and can float on water, 5) They have one valence electron which they readily lose in chemical reactions.

      What is the difference between alkali and alkaline earth metals?

      The main differences are: 1) Alkali metals are in Group 1 of the periodic table, while alkaline earth metals are in Group 2, 2) Alkali metals have one valence electron, while alkaline earth metals have two, 3) Alkali metals are more reactive than alkaline earth metals, 4) Alkali metals form +1 ions, while alkaline earth metals form +2 ions, 5) Alkali metals are softer and have lower melting points than alkaline earth metals.

      Why do alkaline earth metals have similar chemical properties?

      Alkaline earth metals have similar chemical properties because they all have two valence electrons in their outermost shell. This electronic configuration leads to similar behavior in chemical reactions, such as their tendency to form +2 ions and their reactivity with water and oxygen. Additionally, as you move down the group, the atomic radius increases, but the effect of the two valence electrons remains consistent, resulting in gradual trends in properties.

      Prerequisite Topics

      Understanding the foundations of chemistry is crucial when delving into the fascinating world of Group 1 and Group 2 elements, also known as alkali and alkaline earth metals. To fully grasp the properties and behaviors of these important element groups, it's essential to have a solid understanding of two key prerequisite topics: the history and development of the periodic table and atomic structure.

      The periodic table is the cornerstone of modern chemistry, and its evolution provides invaluable context for understanding how elements are organized. By exploring the history and development of the periodic table, students gain insight into how scientists discovered patterns in elemental properties, leading to the grouping of elements with similar characteristics. This knowledge is particularly relevant to Group 1 and Group 2 elements, as their placement in the periodic table reflects their shared properties and reactivity trends.

      Moreover, comprehending the atomic structure is fundamental to understanding why alkali and alkaline earth metals behave the way they do. The electronic configuration of these elements, particularly their valence electrons, plays a crucial role in determining their chemical and physical properties. By mastering concepts related to atomic structure and bonding, students can better explain phenomena such as the high reactivity of Group 1 metals or the tendency of Group 2 elements to form divalent compounds.

      The periodic table groups are not arbitrary divisions; they are the result of careful observation and analysis of elemental properties. Understanding how these groups were established helps students appreciate the logic behind the classification of alkali and alkaline earth metals. It also provides context for why these elements exhibit gradual changes in properties as you move down each group.

      Additionally, a strong foundation in atomic structure and bonding is essential for grasping concepts like ionization energy, atomic radius, and electronegativity all of which are critical in explaining the characteristic behaviors of Group 1 and Group 2 elements. For instance, the low ionization energies of alkali metals can be directly linked to their atomic structure, explaining their high reactivity and tendency to form positive ions.

      By thoroughly understanding these prerequisite topics, students will be well-equipped to explore the intricacies of alkali and alkaline earth metals. They will be able to draw connections between the elements' positions in the periodic table, their electronic structures, and their observed chemical and physical properties. This comprehensive approach not only enhances learning about Group 1 and Group 2 elements but also provides a solid framework for understanding broader concepts in chemistry.

      In this lesson, we will learn:
      • To recall properties of the alkali metals
      • To understand the trend in properties found in the alkali metals
      • To apply knowledge of electronic structure and bonding to explain the trends in group 1 properties.
      • To know the common reactions of group 1 and group 2 elements.

      Notes:
      • We saw earlier that the Periodic Table is arranged, left to right, by proton number and number of outer shell electrons. The number of outer shell electrons dictates the chemical properties of an element.
        Therefore, it is easy to see which elements have similar properties to each other – they will be in the same column of the table as each other, the columns which we call groups.
        Groups 1 and 2 have 1 and 2 electrons in the outer shell, respectively. Their behaviour in chemical reactions is similar because of the similar outer shell configuration.

      • The alkali metals in group 1 and the alkaline earth metals in group 2 are very well-studied groups of elements, with clear patterns in how their properties change.
        They have been known for quite a long time through their compounds, like their metal oxides, because they are reactive and quickly form compounds with oxygen in air. Their oxides all produce alkaline solutions in water, which is how they get their names. Because of their reactivity, they were not isolated as metals until later on.

      • Alkali metals and alkaline earth metals have the following properties:
        • They are relatively soft metals.
        • They are relatively low density metals.
        • They have relatively low melting points compared to metals in general.
        • They are reactive, more so than d-block metals in general and they react vigorously with water.
          • The products of this reaction are hydrogen gas and a metal hydroxide – this forms an alkaline solution, which gives the two groups their names.

      • As you go down the group, the properties of the elements change in the following ways:
        • The metals get softer.
        • The melting point of the metals gets lower.
        • The metals get denser.
        • The metals get more reactive.
        • Ionization energies of the elements decrease.

      • The alkali metals all have a valence of 1 and alkaline earth metals have a valence of 2. As they are metals, they form ionic compounds with non-metals. In these compounds, you’ll see alkali metals with a 1+ charge (such as in NaCl), whereas alkaline earth metals will hold a 2+ charge (such as in MgCl2).

      • As you descend the table, a similar trend in the change in properties is observed in both groups. For example, both group 1 and group 2 show a decrease in ionization energies going down the group.

        As explained in Periodic Trends: Ionization Energy, this is due to greater shielding from more inner electron shells between the nucleus and the outer shell. Greater shielding makes losing the one (in group 1) or two (in group 2) outer shell electrons increasingly easy, and therefore reactivity in general increases going down the two groups.

      • The reactions of group 1 and group 2 metals are very similar because they have similar outer shell configurations and elements in both groups are prone to losing their outer shell electrons.
        • One of the most common reactions with group 1 and 2 elements is the reaction with water. Group 1 and group 2 metals are both known for reacting vigorously with water compared to other metals. The equation can be written generally as:

          For a group 1 metal:
          2M (s) + 2H2O (l)  \, →  \, 2MOH (aq) + H2 (g)


          For a group 2 metal:
          M (s) + 2H2O (l)  \, →  \, M(OH)2 (aq) + H2

          The stoichiometry is slightly different because two hydroxide groups will be bonded to the group 2 metal because it forms a 2+ ion, unlike the 1+ ion of a group 1 metal.


        • Another common reaction is the group 1 or 2 reaction with oxygen. This can be written generally as:

          For a group 1 metal:
          4M (s) + O2 (g)  \, →  \, 2M2O (s)


          For a group 2 metal:
          2M (s) + O2 (g)  \, →  \, 2M2O (s)

          • Again, the stoichiometry is slightly different in the two reactions.
          In the reaction with heavier group two metals, a metal peroxide can form, such as with barium and strontium:
          M (s) + O2 (g)  \, →  \, MO2 (s)

          The peroxide ion (-O-O-) has two oxygen atoms, each with a weak covalent bond to the other, and a 1- charge (O22- overall). It does not form in this reaction with the smaller group 2 metals such as Be or Mg because they form highly polarizing ions.
          These are small with a very high charge density which draws the 2- negative charge in the peroxide ion towards it with great strength (they polarize the ion). The peroxide ion becomes a stable O2- oxide ion and the weak O-O covalent bond breaks.
          The heavier, less dense and less polarizing ions like Sr2+ and Ba2+ cannot do this, so the metal peroxide is stable.

        • The reaction with oxygen to form a metal oxide occurs spontaneously in air and when the metals are heated by flame, showing a distinct colour.

          • The following is a table of the flame colours observed in the group 1 and 2 elements:

          Group 1 elements

          Group 2 elements

          Lithium: red

          Beryllium: white/colourless

          Sodium: orange/yellow

          Magnesium: white/colourless

          Potassium: lilac

          Calcium: red/orange

          Rubidium: red

          Strontium: dark red

          Caesium: blue/violet

          Barium: green


        • The reaction between the group 1 and 2 metals with chlorine can be written generally as:


          • For a group 1 metal:
          2M (s) + Cl2 (g)  \, →  \, 2MCl (s)


          For a group 2 metal:
          M (s) + Cl2 (g)  \, →  \, 2MCl (s)


      • The solubility of many group 1 and 2 metal compounds have trends down the groups. Generally, group 1 metal compounds are more soluble than any group 2 analogue.
        For example, you would predict that potassium hydroxide has greater solubility (in water at a fixed temperature) than calcium hydroxide.
        • Solubility of the metal hydroxides increases down the group. This is true for both group 1 and group 2 hydroxides, and as said above the group 1 hydroxides are much more soluble
        • Solubility of the group 2 metal sulfates decreases down the group. Barium sulfate is the product we look for in the common test for sulfate ions. The white precipitate is obvious as BaSO4 is extremely insoluble.
        • Solubility of the group 2 metal carbonates generally decreases down the group. In group 1, solubility actually increases down the group.
        Explaining the solubility trends is not necessary. In short, it’s because of two enthalpy measurements (lattice enthalpy and enthalpy of hydration) that change at different rates down the groups, depending on the different anions (sulfate, carbonate, etc) used... don’t worry about it!